Acids and Bases, or more specifically, H+ ions and OH- ions always are proportionate within the solution of water. Unfortunately for us and the rest of the people that must deal with the numbers involved, the proportion is not at all linear. the concentration of H+ ions [H+] in any given solution is 10^(-pH) M.

Arrhenius definition of acid/base:
-acids increase H+ or hydronium concentrations in water
-bases increase OH or decrease hydronium concentrations in water

Bronsted-loury definition:
-Acids=proton donors
-Bases=proton acceptors

*All Arrhenius acids/bases are Bronsted-Lowry acids/bases, but not all Bronsted-Lowry acids/bases are Arrhenius acids/bases.

Monoprotic-donates 1 proton
Amphiprotic-can be either Bronsted acid or base
Amphoteric-can be either Lewis acid or base
Polyprotic- donates more than 1 protons (i.e:sulfuric, phosphoric, and oxalic acid)

*adduct = product of acid/base reaction


How to Tell if a Substance is an Acid or a Base:

• Anions that are conjugate bases of strong acids are such weak bases that they have no effect on the pH (it is neutral)
  
  
• Alkali metals and alkaline earth metals cations have no effect on the pH (it is neutral)
  
  
• There are many basic anions - All conjugate acids of a weak acid.
  
  
• Anions from polyprotic acids can be either acids or base depending on the KA
  
  
• Acidic cations will be 2+ or 3+ charges in the transition metals. They act as Lewis Acids and hydrate in water.
• When given something like HCO3 and you have to figure out if it acts as a base or an acid, react it with water as both an acid and a base. Then compare the Ka and the Kb of both reactions. If the Ka>Kb then the solution would be acidic.
  

Formulas:

1. pH= -log [H3O+]
   
   
   
2. pOH= -log [OH-]
   
   
   
3. pH +pOH=14
   
   
   
4. [H3O] = 1x10^-pH
   
   
   
5. [OH-]= 1x10^-pOH
   
   
   
6. Kw= [H30+][OH-]= 1x 10^-14
   
   
   
7. Ka * Kb = Kw

8. Ka= [H30+][A-]/[HA]


9. Kb=[BH+][OH-]/[B]


Henderson-Hasselbach Equation:

pH = pKa + log [A-]/[HA] pKa= -log Ka

pOH = pKb + log [HB+]/[B] pKb = -logKb

This equation is helpful when creating a buffer solution. A buffer causes solutions to resist a change in pH when a strong acid or base is added. If given a specific pH, put the number in for pH and Ka of the weak acid/base and you can determine the ratio of conjugate : weak acid.
The ratio of conj:acid cannot exceed 1M :0.1M. If exceeds, no longer buffer.

In the Henderson-Hasselbach equation, if the ratio of the conjugate to its subsequent acid/base is 1, the log of 1 is 0 so pH=pKa but when ratio>1, pH>pKa and vice versa.

IF KA <10^-5 or so (really small), you can ignore x. So instead of KA= [x^2]/[HB-x], you can write it as KA= [x^2]/[HB].
Ex) What is the pH of a 0.5 M solution of acetic acid. (1.8 x 10^-5= KA at 25 degrees Celsius)
1.8 x 10^-5 =[x^2]/[0.5]
x^2= 0.000009
x= 0.003, which is [H+]
pH= -log [0.003]= 2.522


Lewis Acid/Base
electron acceptor = acid (AAE--acids accept electrons)
electron donor = base



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Acid-Base Reactions

1. Strong Acid + Strong Base
   • net ionic = H30+ + OH- ~> H20
   • very large K value = all product
   • pH depends on excess reactant, if any. the remaining reactant then dissociates into either its H+ or OH- ion which you can then plug into the pH/pOH equation
2. Strong Acid + Weak Base
   • the strong acid dissociates into the hydronium ion
   • the hydronium then reacts with the OH- ions after the weak base undergoes hydrolysis
   • if at the equivalence point, it will form an acidic solution (NOT BECAUSE OF STRONG ACID THOUGH)
   • the pH at the eqivalence point is driven by the conjugate acid of the weak base undergoing hydrolysis and reforming weak base and H+
     • the degree at which the reformation of weak base and H+ happens depends on the Ka of the conjugate acid.

Eg: NH3 + H20 <~> NH4+ + OH- Kb = 1.8 x 10^(-5)
H30+ + OH- <~> 2 H20 K = 1.0 x 10^(-14)
H30+ + NH3 <~> H20 + NH4+ Knet = 1.8 x 10(-9)

3. Weak Acid + Strong Base

• same as Strong Acid + Weak Base but opposite
• Important parts to see on the curve

1. the pH before titration begins
2. the pH at the equivalence point

• driven by the conjugate base of the weak acid undergoing hydrolysis and reforming weak acid and OH-

3. the pH at the halfway point (half-equivalence point)

• at this point, the concentration of weak acid and conjugate base are equal. This makes the ratio in the Hans equation 1 SO pKa = pH
• This is buffer zone

4. Weak Acid + Weak Base

• whichever K is greater determines the pH (if at equivalence)
• double buffer situation: one before equivalence point and one after

Solving the pH of an Acid-Base Solution

1. Calculate the Stoich
2. Calculate the Concentrations
3. Find the pH

*Polyprotic Acids depend primarily on the first acidic hydrogen's Ka
Equivalence Point- point at which moles of acid = moles of base

- pH is not necessarily 7 at equivalence point

strong acid/strong base pH = 7
strong acid/weak base pH < 7
weak acid/strong base pH > 7
weak acid/weak base Depends on Ka and Kb...the larger one influences pH

Buffers are common ion problems. When you have a common ion, it lowers the overall solubility.

Titrations between a weak acid and a weak base are usually not done, because the equivalence point is hard to judge correctly.

Polyprotic Acids:
A titration graph between a diprotic acid with a strong base has a little bump in the curve. This differs from the curves of a strong acid titrated with a strong base and a weak acid titrated with a srong base, a strong acid with a weak base and a weak acid with a weak base. Those curve do not have a extra bump in the curve.

Inductive Effect: extent to which adjacent atoms or groups of atoms attract electrons from another part of a molecule. It comprises the attraction of electrons from adjacent bonds by more electron-negative atoms.


*oxides of nonmetals are acid anhydrides
*oxides of metals are basic anhydrides
*adding water reacts to form acid/base

• the more oxygens an oxoacid has, the stronger the acid [H2SO4 > H2SO3.. HClO4 > HClO3 > HClO2 > HClO] because the electronegative oxygens draw electron density away from the central atom and therefore from the H-O bond, making the substance more polar so the H leaves more easily.
• the more electronegative the central atom, the stronger the oxoacid [H3PO4 < H2SO4 < HClO4] because the more electronegative atom in the center makes the O-H bond more polar.
• binary acid strength depends on the size of the atom [HF < HCl < HBr < HI] because greater distance means weaker attraction.